Download Energy - Haiku for Ignatius

How did we arrive
at this picture of
the atom?

Rutherford’s Atom



Lord Kelvin
The entire atom is like
a positively charged
pudding with
negatively charged
electrons dispersed
throughout like
raisins.
This view lasted until
1910.


Earnest Rutherford
Used a Particle Accelerator to discover that the majority of the
space in an atom is empty.
› The nucleus is a concentrated mass in the center and is made
up of positively charged protons.
› The negative particles, named electrons by JJ Thomson, are
located all around the nucleus.

This view was introduced in 1919.




James Chadwick
Working with
Rutherford, he
discovered that there
was excess mass in the
nucleus that was not
accounted for by the
protons.
He named these
neutral, no-charge
particles neutrons.
BUT!! We still don’t
know what the
electrons are doing!
Electromagnetic Radiation is several types of
energy that exhibit wavelike behavior and
travel at the speed of light in a vacuum.
It was eventually discovered
that electrons behave more
like waves than like particles,
perhaps because of their
incredibly small mass…
9.10938188 × 10-31 kg


Energy travels in the form of waves.
Waves have special properties…
1. Wavelength
2. Frequency
3. Speed
1.
Wavelength – the distance between two
consecutive wave peaks or troughs.
› The symbol for wavelength is λ, the Greek letter lambda.
2.
Frequency – the number of waves (cycles) per
second that pass a given point in space.
› The symbol for frequency is ν, the Greek letter nu.
3.
Speed – All types of electromagnetic radiation travel
at the speed of light (c = 2.9979 x 108 m/s).
› The speed of light is the product of the wavelength and
frequency. c = λ ν
› Wavelength and Frequency have an inverse relationship.

Remember ROIGBIV?

Light travels in waves of particles called
photons which can be viewed as little packets
of energy.
› Photons can travel at different frequencies, and
these frequencies correspond to light of different
colors.

Visible light includes light of all frequencies,
and appears white to us.
› When photons of certain frequencies are isolated,
we see light in the various colors of the rainbow.

Of the visible spectrum, violet light has the most
energy, because it travels with the greatest frequency.

Some light travels with higher energies than we can
see. Examples are listed in order of increasing energy.
1.
2.
3.

Ultraviolet Light
X-rays
Gamma Rays
Some light travels with lower energies than we can see.
Examples are listed in order of decreasing energy.
1.
2.
3.
Infrared
Microwaves
Radio waves (FM, shortwave, AM)
What happens
to wavelength as
frequency
increases?
What happens
to the speed?
What is the
wavelength
range for the
visible light
spectrum?
What is the
frequency of red
and purple light?
When atoms receive
energy from some
source, they become
excited.
Eventually, they release
this energy by emitting
photons.
If the photons have
wavelengths within the
visible light spectrum, we
can see them as light.

As atoms take in energy from the heat of
the flame, they become excited.
› What’s actually happening is the atom absorbs
energy, which allows its electrons to temporarily
jump up levels, and occupy higher energy
positions farther away from the nucleus.

The atom can only hold this extra energy in
for so long before it emits the energy in the
form of energy packets, or photons of light.
› As electrons return to their original positions, the
atom returns to ground state, and it loses its extra
energy by shooting off a photon.

The light signature of each metal has been
determined by the use of a special tool
called a spectroscope.
› The spectroscope splits the light of the visible
spectrum into distinct lines of light that are
separated in a way that our eyes cannot
achieve unaided.

Because of this research, we can identify
the metal ions in an unknown solution by
the color they emit when excited.
Two Scientists are credited
with the discovery of the true
nature of light and energy.
Max Planck & Albert Einstein
discovered that energy is
quantized…in other words,
only certain values are
allowed.

Matter and Energy were considered to
be separate entities.

Matter had particles… energy did not.
› The location of energy in space was
considered indeterminable.

It was thought that matter could absorb
or admit any quantity of energy.

Different wavelengths of light carry different amounts
of energy per photon.
› The energy contained in the photon corresponds to the
change in energy that the atom experiences when an
electron goes from an excited state to a lower state. (NOT
necessarily the ground state.)

Because the atom can exist only in specific energy
states, it can contain only an amount of energy equal
to the amount necessary to attain an energy level, not
any amount in between levels.
› We refer to this property as the quantized nature of energy.
› Quantized – only certain values are allowed.
› Think of stairs compared to a ramp…
The photoelectric effect is the phenomenon in
which electrons are emitted from the surface of a
metal when light strikes it.
Only light higher than a
certain energy will cause
electrons to be emitted.
Einstein won the Nobel
Prize in 1921 for his study
of the photoelectric
effect.
1.
Studies in which the frequency of the light is varied
show that no electrons are emitted by a given metal
below a specific threshold frequency, v0.
2.
For light with frequency lower than the threshold
frequency, no electrons are emitted regardless of
the intensity of the light.
3.
For light with frequency greater than the threshold
frequency, the number of electrons emitted
increases with the intensity of the light.
4.
For light with frequency greater than the threshold
frequency, the kinetic energy of the emitted
electrons increases linearly with the frequency of the
light.
Energy is quantized.
It can occur only in discrete
units called quanta.
Electromagnetic radiation,
which was previously thought to
exhibit only wave properties,
seems to show certain
characteristics of particulate
matter as well.
It was later found that matter
exhibits a dual nature as well.

Electrons, which are considered to be
particles also exhibit wavelike properties.
› Excited ions emit photons of light which have a
frequency which corresponds to the energy
change that occurs when an electron jumps
from an excited state to the ground state.

There is a size requirement for the dual
nature to be easily observed.
› Particles that are too big behave like particles.
› Particles that are too small behave like waves.

The Important Point is that matter and
energy are not distinct.
When white light is dispersed, it displays a continuous spectrum of
visible light, like that seen below. Why are there 2 rainbows?
When atoms are excited, they give off of light (flame test lab), This light
can also be dispersed. The spectrum that is produced, while not
continuous like a rainbow, is unique to that element.

Continuous Spectrum
› a spectrum that contains ALL the
wavelengths of visible light.
› All the lines are present.

Line Spectrum
› a spectrum that contains only certain
wavelengths of visible light.
› Only some of the lines of are present.

Hydrogen is a good place to start because
it only has one electron.
› Therefore, all the light it emits comes from the
same source.

Hydrogen emits light with a four-line visible
light spectrum.
› The wavelengths of the light are 410 nm, 434 nm,
486 nm, & 656 nm.

The significance of the existence of atomic
spectrum of hydrogen is that it indicates
that only certain energies are allowed for
the electron in the hydrogen atom.
Pages
287 –288
Niels Bohr (1885 – 1962) was a
Danish physicist.
He developed a quantum
model for the Hydrogen atom.
He proposed that the electron
in the hydrogen atom moves
around the nucleus only in
certain allowed circular orbits.
In order for an electron to
move between levels, it must
gain or lose a specific quantity
of energy.
The wavelengths predicted matched those
that were measured.
 The model correctly fits the quantized
energy levels of the hydrogen atom and
postulates only certain allowed circular
orbits for the electrons.
 Electrons require energy to exist in orbits
farther away from the nucleus because of
the attraction of opposite charges.
 As the electron is brought closer to the
nucleus, (like when it jumps from an excited
state to ground state) energy is released
from the system in the form of a photon..








An electric current is sent through a tube containing
H2(g).
The energy causes the molecules to separate and fly
around in a high energy state.
The atoms collide and recombine, excited electrons
returning to lower energy states & releasing light
energy.
The light is split by a diffraction grating into the line
spectrum of hydrogen.
The distance to the lines in the line spectrum is
measured using a spectroscope.
The wavelengths of the corresponding photons given
off when the electrons release energy are
determined.
The wavelengths are used to determine which
quantum levels the electrons jump from in order to
give off each respective color line in the spectrum.
Mid 1920s
Werner Heisenberg (1901-1976)
Louis de Broglie (1892-1987)
Erwin Schrodinger (1887-1061)
Compares the electron bound to
the nucleus to a standing wave.



Bohr’s model was known to be incorrect by the mid
1920’s.
The motion of a large particle (like a billiard ball) is
predictable.
Electrons, however, do not move predictably.
› De Broglie proposed that the smaller a particle is, the more
it behaves like a wave (light), and the less it behaves like a
particle (billiard ball).

The unpredictable position of an electron is the main
reason why scientists knew that Bohr’s model was
incorrect.
› According to his model, we should know exactly where
the electrons are…but we can’t find them!

But how then, can we determine the position of a
particle, if we do not understand how it moves?

De Broglie and Schrodinger came up with a
VERY complex mathematical equation that
was able to predict the location of an electron
based on the wave-particle theory.

The result is our modern concept of the atom
containing not only orbits, but orbitals.
› Orbital – The space around the nucleus where an
electron is MOST LIKELY TO BE.

There are four types of orbitals and each has its
own specific energy, shape, & # of electrons it
can hold.
› s, p, d, & f

It is important to note that the concept of an
orbital is not a method for finding an exact
location of an electron.

Orbitals are just a prediction based on statistics
and probability. (See next slide)

Electrons do not follow a direct and
predictable path. (the firefly experiment)

Atoms have 2 subdivisions for housing
electrons.
› Principal energy levels (like orbits)
› Sublevels (the orbitals within each orbit)
It is somewhere in its orbital, following an
indeterminable path, in a position which
can only be determined by probability!
 The definition most often used by
chemists to describe the size of the
hydrogen 1s orbital is the radius of the
sphere that encloses 90% of the total
electron probability.


Principle Energy Levels are represented by the
numbers 1-7.

The total number of principle energy levels an
atom has corresponds directly to the row in
which the atom is found on the periodic table.
› For example, Silver, which is found in row 5 has five
principle energy levels (1,2,3,4 & 5).

The higher in number a principle energy level is,
the farther away it is from the nucleus.
› Which means an electron must have higher energy
in order to be there.

There are four sublevels.
› s, p, d, f

Each principle energy level can contain a
number of sublevels (orbitals) equal to or
less than its principle energy level number
up to a maximum of 4.
› PEL 1 can contain only an s orbital.
› PEL 3 contains s, p & d.
› PEL 6 contains s, p, d & f.

Only one of each sublevel is contained in
each principal level, and once these are
full, electrons must move up to the next
energy level.
The s Orbital as a
spherical shape.
 The larger the value of n,
the larger the sphere.
 In the picture, 1s and 2s are shown as
white rings.

› The blue center represents the nucleus.
› The pink area between the 1s and 2s orbitals
is an area of zero probability where no
electrons exist.




The p Orbital has two lobes separated by a node
at the nucleus.
The p Orbital is only possible for atoms with
principle energy levels ≥ 2.
The larger the value of n, the larger the lobes.
There are three possible orientations in space.
› 2px
› 2py
› 2pz
Which p orbital is pictured? 
The d Orbital has an even more complex shape
than p.
 The d Orbital is only possible for atoms with
principle energy levels ≥ 3.
 The larger the value of n, the larger the shapes.
 There are five possible orientations in space.
› 3dxz
› 3dyz
› 3dxy
› 3dx2-y2
› 3dz2





The f Orbital has the most complex shape.
The f Orbital is only possible for atoms with principle
energy levels ≥ 4.
The larger the value of n, the larger the shapes.
There are seven possible orientations in space.
› 4fz3 – (3/5) zr2
› 4fx3 – (3/5) xr2
› 4fy3 – (3/5) yr2
› 4fxyz
› 4fy(x2 – z2)
› 4fx(z2-y2)
› 4fz(x2 –y2)

Pauli Exclusion Principle states that an atomic
orbital can hold a maximum of two electrons, and
those two electrons must have opposite spins.

Electrons spin much like tops.
› They can spin in one of two directions, clockwise and
counterclockwise.
› We represent the fact that they spin in different
directions by using arrows: ↑ & ↓

It is not important to know which way an electron
is spinning, just that if there are two in one orbital,
they MUST have opposite spin.
n
Orbitals
Orientations
Total # of Electrons
1
s
1
2
2
s, p
1, 3
2+6=8
3
s, p, d
1, 3, 5
2 + 6 + 10 = 18
4
s, p, d, f
1, 3, 5, 7
2 + 6 + 10 + 14 = 32
1s2
2s2, 2p6
3s2, 3p6, 3d10
4s2, 4p6, 4d10, 4f14

Electron Configuration – electron arrangement.

There are two ways to represent electron
configuration:
› Electron configuration notation – Uses the principal
energy level, the orbital shape, and the number of
electrons present.
 Hydrogen: 1s1
› Orbital Diagram – Uses the principal energy level and
the orbital shape plus a series of boxes labeled with
arrows to represent electrons present as well as their
spins.
 Hydrogen:
1s
↑

As atoms increase in # of protons and
electrons, it can become confusing as to
where the additional electrons go.
› The electrons fill up atomic orbitals in a
specific order.

The periodic table of elements can be
used to determine this order.
› Lets consider the first 18 elements on the
periodic table…
› http://www.dayah.com/periodic/

Orbitals fill up in the following order:
› 1s
› 2s, 2p
› 3s, 3p
› 4s, 3d, 4p
› 5s, 4d, 5p
› 6s, 4f, 5d, 6p
› 7s, 5f, 6d, 7p

Do you notice any patterns?
Energy
Order of increasing energy is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f,
5d, 6p, 7s, 5f, 6d, 7p, 8s,
Pages
303 - 307
Electrons in the outer most principle
quantum level of an atom.
 These are the most important electrons to
chemists because these are the electrons
that are involved in bonding.
 The elements in the same group (vertical
column of the periodic table) have the
same valence electron configuration.

› This is why Mendeleev was able to align the
elements by their properties; because elements
with = # of valence electrons behave in
chemically similar ways.

The quantum mechanical model can be used
to explain the arrangement of the elements in
the periodic table.

Quantum mechanics allows us to understand
that similar chemistry exhibited by the
members of a given group arises from the fact
that they all have the same valence electron
configuration.

Only the principle quantum number changes
going down a group on the periodic table.
The Wave Mechanical Model is further
supported by several properties, which are
directly related to electron configuration,
that demonstrate a periodic repetition on
the Table of Elements.
 We will talk about four:

› Metals, Nonmetals, & Metalloids
› Chemical Reactivity
› Ionization Energies
› Atomic Size

Metals
› Metals are ductile and malleable, can conduct heat and
electricity, and have a shiny luster.
› Metals give up electrons to form positive ions.
› The majority of the periodic table, from the staircase to the
left.

Nonmetals
› Do not have the properties of metals.
› Nonmetals gain electrons to form negative ions.
› The upper right corner of the periodic table to the right of
the staircase.

Metalloids
› Exhibit the properties of both metals and nonmetals.
› Include only 7 elements: Si, Ge, As, Sb, Te, Po, & At

Electronegativity is the property of an
atomic nucleus to attract electrons.
› Electronegativity increases going up a group.
› Electronegativity increases going across a row.

The chemical reactivity of an element
depends on its ability to gain or lose
electrons.
› We will look at chemical reactivity through two
properties.
 The property of metals to lose electrons.
 The property of nonmetals to gain electrons.

In the Alkali Metal group, Cesium is much
more likely to lose an electron than Lithium.
› Each has an s1 electron in its valence shell.
› For Li, that electron in principle energy level 2,
and very close to the + charge of the nucleus.
› For Cs, that electron is in principle energy level 7,
and is much further away from the pull of the
positive nucleus.
The same trend is seen in the Alkaline Earth
Metals.
 The most chemically active metals reside in
the lower left hand corner of the table.


In the Halogen group, Fluorine has a much
higher electronegativity than Astatine.
› At has 76 more electrons than F
› The more massive the electron cloud, the bigger
the force of negative repulsion a new electron
must face.
› In F, the nucleus does not have so many
electrons to pull through, and it can attract that
extra electron more easily.


The same trend is seen in Oxygen’s group.
The most chemically active nonmetals are
found in the upper right hand corner.

Ionization Energy – the energy required to
remove an electron from an individual atom in
the gas phase.
› M(g)  M+(g) + eionization
energy

Metals have relatively low ionization energies.
› Metals increase in reactivity going down the group.

Nonmetals have relatively high ionization
energies.
› Nonmetals would rather gain than lose electrons.

Ionization Energy increases going up a group
and going across a row.

Atomic Size in a group
› As we go down a group, we add a principle energy
level of electrons.
› It is easy to see why this would cause an increase in
atomic radius.

As we go across a row, however, atomic radius
decreases, even though we are adding
electrons.
› This is because we are also adding protons.
› As the # of protons increases, so does the nucleus’s
ability to pull the electron cloud in.

Atomic radius increases going down a group,
and decreases going across a collumn.
1.
2.
3.
It is the number and type of valence electrons that primarily
determine an atom’s chemistry.
The organization of the periodic table can be used to
approximate the electron configuration of quantum
mechanical model atoms.
Certain groups have special names that should be known.
•
•
•
•
•
•
•
Alkali Metals
Alkaline Earth Metals
Halogens
Noble Gases
Lanthanides
Actinides
Transition Metals.
4.
Elements can also be divided into Metals and Nonmetals.
•
•
•
Metals give up electrons to form positive ions, and have low
ionization energies. They are found on the left side of the table.
The most chemically reactive metals are in the bottom left
corner of the table. Metals are usually reducing agents.
Nonmetals gain electrons to form negative ions when reacting
with a metal. Nonmetals have high ionization energies. They are
on the right side of the table. The most reactive nonmetals are in
the upper right corner (ignoring the noble gases because they
are inert). Nonmetals are usually oxidizing agents.
The division between metals and nonmetals is defined by the
stair step line. But the division is only approximate. Many of the
elements touching the division line have the properties of both
metals and nonmetals. These elements are called metalloids.







Where are the nonmetals found on the
periodic table?
Where are the metalloids found?
Where are the metals found?
Where are the most chemically active
metals located?
Where are the most chemically active
nonmetals located?
How does the ionization energy of an atom
change periodically?
How does atomic size change periodically?